Aim: How do we perfect our knowledge of the periodic table? Do now: How is the periodic table arranged?

Announcements Flame Test Lab - Thu, Sept. 8 Exam #1- September 12th Free Response # 2 - due Thu, Sep 15th Reading Guide #1 and #2 - due September 12th

Organizing the Elements A few elements, such as gold and copper, have been known for thousands of years - since ancient times Yet, only about 13 had been identified by the year 1700. As more were discovered, chemists realized they needed a way to organize the elements.

Organizing the Elements Chemists used the properties of elements to sort them into groups. In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties One element in each triad had properties intermediate of the other two elements

Mendeleev’s Periodic Table By the mid-1800s, about 70 elements were known to exist Dmitri Mendeleev – a Russian chemist and teacher Arranged elements in order of increasing atomic mass Thus, the first “Periodic Table”

He left blanks for yet undiscovered elements Mendeleev He left blanks for yet undiscovered elements When they were discovered, he had made good predictions But, there were problems: Such as Co and Ni; Ar and K; Te and I

A better arrangement In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number The arrangement used today

Another possibility: Spiral Periodic Table

Trends in Atomic Size First problem: Where do you start measuring from? The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time.

Atomic Size } Radius Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

ALL Periodic Table Trends Influenced by three factors: 1. Energy Level Higher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) More charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect (blocking effect?)

#1. Atomic Size - Group trends H As we increase the atomic number (or go down a group). . . each atom has another energy level, so the atoms get bigger. Li Na K Rb

#1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. Electrons are in the same energy level. But, there is more nuclear charge. Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

Rb K Period 2 Na Li Atomic Radius (pm) Kr Ar Ne H 3 10 Atomic Number

#2. Trends in Ionization Energy Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom). Removing one electron makes a 1+ ion. The energy required to remove only the first electron is called the first ionization energy.

Ionization Energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE.

Symbol First Second Third HHeLiBeBCNO F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

What factors determine IE The greater the nuclear charge, the greater IE. Greater distance from nucleus decreases IE Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. Shielding effect

Shielding The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. Second electron has same shielding, if it is in the same period

Ionization Energy - Group trends As you go down a group, the first IE decreases because... The electron is further away from the attraction of the nucleus, and There is more shielding.

Ionization Energy - Period trends All the atoms in the same period have the same energy level. Same shielding. But, increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 full orbitals.

Trends in Ionic Size: Cations Cations form by losing electrons. Cations are smaller than the atom they came from – not only do they lose electrons, they lose an entire energy level. Metals form cations. Cations of representative elements have the noble gas configuration before them.

Ionic size: Anions Anions form by gaining electrons. Anions are bigger than the atom they came from – have the same energy level, but a greater area the nuclear charge needs to cover Nonmetals form anions. Anions of representative elements have the noble gas configuration after them.

Trends in Electronegativity Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. They share the electron, but how equally do they share it? An element with a big electronegativity means it pulls the electron towards itself strongly!

Electronegativity Group Trend The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has. Thus, more willing to share. Low electronegativity.

Electronegativity Period Trend Metals are at the left of the table. They let their electrons go easily Thus, low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away from others High electronegativity.

Electron Affinity The amount of energy needed to remove an electron from a negatively charged ion. It is also energy released when a single electron is combined with an isolated atom.